The Dissolution Process

Core Concepts

The three energetic steps

For an ionic compound like NaCl in water:

1. Break the ionic lattice — overcome the electrostatic attractions holding Na⁺ and Cl⁻ in the crystal. Energy cost = lattice energy (endothermic).
2. Create a cavity in the solvent — separate water molecules from each other (endothermic, usually smaller).
3. Hydrate the ions — water's partially negative oxygen end surrounds Na⁺, its partially positive hydrogen ends surround Cl⁻; the ion–dipole attractions release energy (exothermic).

ΔH_soln = ΔH_lattice + ΔH_hydration

(with hydration negative). For NaCl, breaking the lattice costs slightly more than hydration releases, so ΔH_soln = +3.9 kJ/mol — mildly endothermic, yet salt dissolves readily. The driving force is entropy: the crystal's few ordered ions become countless randomly dispersed hydrated ions.

Endothermic and exothermic dissolution

• Endothermic (ΔH_soln > 0): the solution feels cold because it absorbs heat. Examples: NH₄NO₃ (+25.7 kJ/mol — instant cold packs) and KCl (+17.2 kJ/mol); both dissolve anyway because entropy wins.
• Exothermic (ΔH_soln < 0): the solution warms up. Examples: NaOH (−44.5 kJ/mol) and CaCl₂ (−81.7 kJ/mol — used in hot packs).
• The sign of ΔH_soln tells you about heat, not about whether dissolution happens.

Like dissolves like

Mixing is favorable when the solute–solvent interactions are similar in kind and strength to the interactions being broken:

• Polar solvent + ionic/polar solute: water dissolves NaCl (ion–dipole) and ethanol (hydrogen bonds).
• Nonpolar solvent + nonpolar solute: hexane dissolves wax and cooking oil (dispersion forces).
• Polar + nonpolar: poor mixing — water and oil separate because replacing strong water–water hydrogen bonds with weak water–oil contacts costs more than the entropy gain provides.

"Like dissolves like" is a rule of thumb about intermolecular force matching, not a law — some polar solutes are only slightly soluble, and molecular size matters too.

Factors that speed up dissolution

Thermodynamics says whether dissolution can happen; kinetics says how fast. Stirring brings fresh solvent to the surface and carries solute away; heat gives particles more kinetic energy (and, for most solids, raises solubility — topic 3); smaller particles expose more surface area per gram.

Saturated, unsaturated, and supersaturated solutions

• Unsaturated: can dissolve more solute at that temperature.
• Saturated: holds the maximum solute at that temperature; dissolved and undissolved solute are in dynamic equilibrium (dissolving and precipitating at equal rates).
• Supersaturated: temporarily holds more than the saturation limit (made by cooling a hot saturated solution carefully). It is unstable — a seed crystal or disturbance triggers rapid precipitation; this underlies "hot ice" (sodium acetate) demonstrations and honey crystallization.

How It Works / Step-by-Step Process

Worked example 1: energy bookkeeping for NaCl

Problem. For NaCl, breaking the lattice requires +788 kJ/mol, hydration of the ions releases −784 kJ/mol (summed over both ions). Find ΔH_soln and classify the process.

Solution.

1. Write the formula for the enthalpy of solution: ΔH_soln = ΔH_lattice + ΔH_hydration
2. Substitute: ΔH_soln = (+788 kJ/mol) + (-784 kJ/mol) = +4 kJ/mol (Measured: +3.9 kJ/mol.) Units: kJ/mol — energy per mole of dissolved salt ✓
3. Positive sign → endothermic (the solution cools imperceptibly). Yet NaCl dissolves because the entropy of dispersing the ions is large and positive — spontaneous despite the small energy cost.

Worked example 2: predicting "like dissolves like" outcomes

Problem. Predict which pairs mix well: (a) NaCl in water or hexane; (b) cooking oil in hexane or water; (c) ethanol in water.

Solution.

1. (a) NaCl is ionic → needs ion–dipole forces with a polar solvent → dissolves in water; hexane offers only weak dispersion forces, so the lattice energy is not repaid → insoluble.
2. (b) Oil is nonpolar → dispersion forces only → mixes with hexane; water cannot repay oil's dispersion attractions → two layers.
3. (c) Ethanol has an –OH group → hydrogen bonds with water; its small alkyl tail is tolerated → miscible with water in all proportions.

Worked example 3: cold pack reasoning

Problem. A cold pack contains solid ammonium nitrate and a pouch of water. Squeezing the pack mixes them, and the pack gets cold. Explain using ΔH_soln and entropy.

Solution.

1. NH₄NO₃ dissolution is endothermic: ΔH_soln = +25.7 kJ/mol — lattice and cavity costs exceed the hydration energy released.
2. The reaction draws that energy from its surroundings (the pack and your skin), which is why it feels cold.
3. Spontaneity comes from entropy: one ordered crystal becomes many randomly dispersed NH₄⁺ and NO₃⁻ ions — a large ΔS that outweighs the unfavorable enthalpy. "Spontaneous" does not mean "exothermic."

Common Confusions